Chemical Mechanisms
The corrosion cell For corrosion to occur, the formation of a corrosion cell must be considered. The corrosion cell contains four basic components: the anode, cathode, electrolyte and metallic path.12 The anode is generally the site of corrosion, where the metal atoms go into the solution as metal ions through the process of oxidation. The cathode is the site of reduction, and consumes the electrons generated from the reaction in the anode. The electrolyte is a solution that is electrically conductive. Lastly, the metallic path describes the ionic flow as well as the electron flow.1112 In an iron corrosion cell, the anode is the iron-based metal (surface), the cathode doesn’t have a physical surface but it is effectively the electrolyte, which could be a solution of water. The electrons flow from the iron surface to the solution, and the ion (positive charge current) flows likewise. The negative ions (OH-) flows from the solution to the iron surface, but it is quickly precipitates as Fe(OH)2 before reaching the iron.12 A removal of any of the components will effectively cease the process of corrosion. Formation of rust For the process of corrosion of iron, there can be many simultaneous chemical mechanisms occurring at any given time. Figure 2 is an example. By oxidation, Fe2+ ions are released. In a base, the Fe2+ reacts with OH- ions to form Fe(OH)2, commonly known as rust.12 Reaction (1) is an oxidation reaction, and is also called the “anodic reaction,” since this occurs in the anode.10 It is noted, however, that the ferrous ions linger around the metal surface of iron. The oxidation reaction generates ferrous ions and electrons. When iron is oxidized, it is transitioning in a higher valence state.12 In the presence of oxygen and water, Fe2+ ions can be further oxidized into Fe3+.13 If the process of iron oxidation is stopped, then the process of corrosion will also be effectively terminated. By reduction, OH- will be generated in a neutral or alkaline solution through the process described in equation (3) to (6), depending on the oxygen concentration in the solution, and the pH. Any of these reactions (depending on the conditions) is then called the “cathodic reaction.”10 The electrons from the anodic reaction are consumed here and in the reduction reaction, there is a decrease in valence state.12 A “high pH” could mean that the solution is neutral or basic (i.e. not acidic). In many sources, reaction (6) is typically described since many natural solutions (like water and seawater) either have a neutral or alkaline pH. This means, however, that the corrosion reaction is very dependent on the availability of oxygen. The ferrous ions are deposited into the water solution as the iron is being oxidized, and the hydroxide ions likewise deposited into the water solution. In the solution, reaction (7) occurs to form the Fe(OH)2 precipitate. This is not soluble in water, and it usually forms on the iron surface.12 The hydroxide ions in reaction (7) can come from reaction (5) or (6), or through the hydrolysis of water. Moreover, iron (II) hydroxide, the product of reaction (7), can be further reduced into iron (III) hydroxide in the presence of water and oxygen. With the availability of oxygen, iron (II) hydroxide (Fe(OH)2) hydrates to form iron (III) oxide-hydroxide (Fe2O3'·'''H2O).12 'Side reactions' The formation of Fe(OH)2 through the process detailed above is just one of the many simultaneous reactions occurring in the corrosion cell. Details of reactions in the corrosion of iron include the below mechanisms.12 Iron, in the form of Fe3+, will precipitate as Fe2O3, which is independent of the potential since it is a chemical process that does not depend on electron availability in the solution.13 Equilibrium thermodynamics and kinetics 'Factors affecting corrosion' *The process of corrosion is highly dependent on the concentration of oxygen in the solution. **The corrosion of iron is very dependent on the availability of oxygen, as described in reaction (6) and consequently, reaction (7). Thus, for any temperature, the corrosion rate increases with oxygen concentration. **At higher temperatures, the solubility of oxygen is lower.11 *Corrosion is also affected greatly by the temperature. Generally, as the temperature increases, so does the rate of corrosion. This is due to the mixture of a lower activation energy (since temperature is higher) for the chemical and electrochemical reactions, and an increased rate of diffusion in the electrolyte. 11 The acidity of the solution (water solution) is also a determining factor in the corrosion rate. A lower pH (i.e. more acidic) solution will tend to increase the rate of corrosion.10 Reaction (3) occurs in a solution that is acidic, and with more H+ ions available, this cathodic reaction will consume the electrons generated from the anodic reaction (1). 'Thermodynamics and kinetics''' The rate of corrosion is very dependent on the oxidation and reduction reactions at the anode and cathode, respectively. Thus, any change in these “redox” reactions will affect the corrosion of iron. Moreover, since the oxidation and reduction reactions are simultaneous reactions occurring, the oxidation proceeds as quickly as the reduction of the cathode. 10 The anodic reaction of iron atom to iron ions, as exemplified in reaction (1), is basically the corrosion reaction. It governs the whole corrosion process, and without it, the corrosion mechanism is terminated. However, reaction (1) is not a spontaneous reaction. The ability of a metal to be corroded depends on the reduction or oxidation potential of the metal. 10 A metal with a lower potential is more readily corroded than that with a higher potential. Table 1 displays the potentials of metals compared hydrogen, and is measured in voltage (V). Because it is not possible to determine the value of the difference in potential (∆V) between an electrode and the solution experimentally, a reference electrode of 0 V must be introduced, and all other electrodes are referred to this.10 In Table 1, the reference electrode here is hydrogen. Metals with a lower standard, or electrode, potential tend to corrode more than metals with higher electrode potential. Iron has an electrode potential of -0.44 V, and thus is has the tendency to be the anode and corrode, with respect to hydrogen. The electrode potential must be achieved before corrosion occurs. The cathodic reaction is also a very significant because this reaction influences the corrosion rate at steady state. The consumption of electrons (from the oxidation or anodic reaction) causes the higher anodic dissolution rate.10 For reaction (1), the electrons in the products are used up by the cathodic reaction; therefore, as stated by Le Chatelier's principle, since the product is removed, the equilibrium shifts to the right, and the rate of the anodic reaction is higher.ise, for example in reaction (6), there is a decrease in oxygen supply, then the cathodic reaction is interrupted, and there would be a decrease in the rate of anodic dissolution (which then decreases the rate of corrosion).10